CHEM101: General Chemistry I

One useful way of organizing our understanding of matter is to think of a hierarchy that extends down from the most general and complex to the simplest and most fundamental. The orange-colored boxes represent the central realm of chemistry, which deals ultimately with specific chemical substances, but as a practical matter, chemical science extends both above and below this region.

Alternatively, when we are thinking about specific samples of matter, it may be more useful to re-cast our classification in two dimensions:

Notice, in the bottom line of boxes above, that mixtures and pure substances can fall into either the homogeneous or heterogeneous categories.

Since this latter schema covers virtually all matter that we encounter in the world, it is important that you thoroughly understand the meaning of each part of it. We will begin by looking at the distinction represented in the top line of the diagram.

Homogeneous and Heterogeneous: It is a Matter of Phases

Homogeneous matter (from the Greek homo = same) can be thought of as being uniform and continuous, whereas heterogeneous matter (hetero = different) implies non-uniformity and discontinuity. To take this further, we first need to define uniformity in a more precise way, and this takes us to the concept of phases.

A volume of water, a chunk of ice, a grain of sand, a piece of copper – each of these constitutes a single phase, and by the above definition, is said to be homogeneous.

A sample of matter can contain more than a single phase; a cool drink with ice floating in it consists of at least two phases, the liquid and the ice. If it is a carbonated beverage, you can probably see gas bubbles in it that make up a third phase.

The phase boundaries in an oil-vinegar salad dressing are easy to see...

...but those in milk are so small that they are only revealed by scattered light.

Phase Boundaries

Each phase in a multiphase system is separated from its neighbors by a phase boundary, a thin region in which the intensive properties change discontinuously. Have you ever wondered why you can easily see the ice floating in a glass of water although both the water and the ice are transparent? The answer is that when light crosses a phase boundary, its direction of travel is slightly bent, and a portion of the light gets reflected back; it is these reflected and distorted light rays emerging from that reveal the chunks of ice floating in the liquid.

If, instead of visible chunks of material, the second phase is broken into tiny particles, the light rays usually bounce off the surfaces of many of these particles in random directions before they emerge from the medium and are detected by the eye. This phenomenon, known as scattering, gives multiphase systems of this kind a cloudy appearance, rendering them translucent instead of transparent.

Two very common examples are ordinary fog, in which water droplets are suspended in the air, and milk, which consists of butterfat globules suspended in an aqueous solution.

Getting back to our classification, we can say that:

Dichotomies (either-or classifications) often tend to break down when closely examined, and the distinction between homogeneous and heterogeneous matter is a good example; this is really a matter of degree, since at the microscopic level all matter is made up of atoms or molecules separated by empty space! For most practical purposes, we consider matter as homogeneous when any discontinuities it contains are too small to affect its visual appearance.

How large must a molecule or an agglomeration of molecules be before it begins to exhibit properties of a being a separate phase? Such particles span the gap between the micro and macro worlds, and have been known as colloids since they began to be studied around 1900. But with the development of nanotechnology in the 1990s, this distinction has become even more fuzzy.

Pure Substances and Mixtures

The air around us, most of the liquids and solids we encounter, and all too much of the water we drink consists not of pure substances, but of mixtures. You probably have a general idea of what a mixture is, and how it differs from a pure substance; what is the scientific criterion for making this distinction?

To a chemist, a pure substance usually refers to a sample of matter that has a distinct set of properties that are common to all other samples of that substance. A good example would be ordinary salt, sodium chloride. No matter what its source (from a mine, evaporated from seawater, or made in the laboratory), all samples of this substance, once they have been purified, possess the same unique set of properties.

A mixture, in contrast, is composed of two or more substances, and it can exhibit a wide range of properties depending on the relative amounts of the components present in the mixture. For example, you can dissolve up to 357 g of salt in one litre of water at room temperature, making possible an infinite variety of salt water solutions. For each of these concentrations, properties such as the density, boiling and freezing points, and the vapor pressure of the resulting solution will be different.

Is Anything Really Pure?

Ivory soap used the marketing tagline, "99 44/100% Pure: It Floats"

This description of Ivory Soap is a classic example of junk science from the 19th century. Not only is the term pure meaningless when applied to an undefined mixture such as hand soap, but the implication that its ability to float is evidence of this purity is deceptive. The low density is achieved by beating air bubbles into it, actually reducing the purity of the product and, by diluting the soap with air bubbles, cheating the consumer.

Similarly, those of us who enjoy peanut butter would never willingly purchase a brand advertised as impure. But a Consumer Reports article published some years ago showed a table listing the number of "mouse droppings" and "insect parts" (presumably from peanut storage silos) they found in samples of all the major brands. Bon Appetit!

Finally, we all prefer to drink pure water, but we don't usually concern ourselves with the dissolved atmospheric gases and ions such as Ca²⁺ and HCO₃^– that are present in most drinking waters. But these harmless impurities and are always present in those pure spring waters.

However, these and similar substances could seriously interfere with certain uses to which we put water in the laboratory, were we customarily use distilled or deionized water.

Even these still contain some dissolved gases and occasionally some silica, but their small amounts and relative inertness make these impurities insignificant for most purposes. When water of the highest obtainable purity is required for certain types of exacting measurements, it is commonly filtered, de-ionized, and triple-vacuum distilled.

But even this chemically pure water is a mixture of isotopic species: there are two stable isotopes of both hydrogen (H¹ and H², often denoted by D) and oxygen (O¹⁶ and O¹⁸) which give rise to combinations such as H₂O¹⁸, HDO¹⁶, etc., all of which are readily identifiable in the infrared spectra of water vapor.

Interestingly, the ratio of O¹⁸/O¹⁶ in water varies enough from place to place that it is now possible to determine the source of a particular water sample with some precision. And to top this off, the two hydrogen atoms in water contain protons whose magnetic moments can be parallel or antiparallel, giving rise to ortho- and para-water, respectively.

Operational and Conceptual Classifications of Matter

Since chemistry is an experimental science, we need a set of experimental criteria for placing a given sample of matter in one of these categories. There is no single experiment that will always succeed in unambiguously deciding this kind of question. However, there is one principle that will always work in theory, if not in practice. This is based on the fact that the various components of a mixture can, in principle, always be separated into pure substances.

Consider a heterogeneous mixture of salt water and sand. The sand can be separated from the salt water by the mechanical process of filtration.

Similarly, the butterfat contained in milk may be separated from the water by a process called churning, in which mechanical agitation forces the butterfat droplets to coalesce into the solid mass we know as butter.

These examples illustrate the general principle that:

Turning this around, we have an operational definition of heterogeneous matter: if, by some mechanical operation we can separate a sample of matter into two or more other kinds of matter, then our original sample was heterogeneous.

To find a similar operational definition for homogeneous mixtures, consider how we might separate the two components of a solution of salt water. The most obvious way would be to evaporate off the water, leaving the salt as a solid residue. Thus a homogeneous mixture can be separated into pure substances by undergoing appropriate partial changes of state – that is, by evaporation, freezing, etc.

Note the term partial in the above sentence; in the last example, we evaporate only the water, not the salt (which would be very difficult to do anyway!) The idea is that one component of the mixture is preferentially affected by the process we are carrying out. This principle will be emphasized in the following examples.

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Source: Stephen Lower, http://www.chem1.com/acad/webtext/pre/pre-1.html#A2A
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